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EDEXCEL Chemistry a level > CHEMICAL BONDING AND STRUCTURE > Flashcards

CHEMICAL BONDING AND STRUCTURE Flashcards

Giant structures Discrete molecules Physical properties (88 cards)

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1
Q

metallic bonding

A

strong electrostatic attraction between the nuclei of metal cations and delocalised electrons.

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2
Q

metals properties

A 
high mp/bp
high thermal/electrical conductivity
malleable
sonorous
ductile
low IE's
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3
Q

Why do metals have high mp’s?

A

Need to break strong forces of attraction between the cation nuclei and delocalised electrons.
also giant lattice structure requires much energy to break

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4
Q

What determines bp/mp’s?

A

no. of delocalised electrons (group 1 is low while d-block metals are higher)
size of cation (smaller radii has closer electrons to nuclei higher mp)

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5
Q

electrical conductivity in metals

A

movement of delocalised electrons towards the positive terminal of a cell with a potential difference.

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6
Q

thermal conductivity

A

delocalised electrons passing KE along metal.

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7
Q

malleability

A

shape configuration

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8
Q

ductility

A

ability to draw metal into a wire

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9
Q

what does ductility and malleability depend on?

A

the cation and electron movement ability. layers slide over one another and delocalised electrons prevent strong forces of repulsion between cations.

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10
Q

ionic bonding

A

strong electrostatic attraction between oppositely charged ions.

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11
Q

arrangement of ionic substances

A

giant ionic lattice in which electrostatic attraction occurs in all directions.

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12
Q

how do you determine the strength of electrostatic attractions in an ionic substance?

A

calculation of energy per mole of solid needed to separate ions to infinity in which ions can no longer interact

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13
Q

how is strength of ionic bonding affected by size of cation?

A

as cation size increases, amount of energy needed to separate them decreases

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14
Q

how is strength of ionic bonding affected by size of anion?

A

increases as anion size increases

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15
Q

how does size of ions affect energy needed to separate ionic substances to infinity?

A

smaller substances require more energy to separate

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16
Q

why must data of ionic radii all come from the same source?

A

because radii is difficult to measure and several methods could conflict in accuracy.

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17
Q

properties of ionic substances

A

high mp/bp
brittle
water-soluble
poor electrical conductivity as a solid

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18
Q

why do ionic substances have high melting points?

A

giant lattice networks of oppositely charged ions have combined large electrostatic forces.

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19
Q

why are ionic substances brittle?

A

as stress causes ionic layers to slide over one another, meaning same charge ions are next to each other, repelling

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20
Q

why are ionic substances water soluble

A

polarity of water molecues allows separation of ions

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21
Q

reason for electrical conductivity when molten

A

previously fixed electrons become mobile, allowing a migration of electrons towards the positive terminal

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22
Q

what sort of current is used in electrolysis

A

direct

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23
Q

covalent bonding

A

2 atoms overlapping atomic orbitals of which contain a singular electron

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24
Q

3 ways covalent bonds interact

A

sigma (2 s orbitals)
sigma (2 P orbitals)
pi (2 P orbitals)

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25
how may pi orbitals form? how does this affect the molecule?
sigma bonds may lead to the formation of pi bonds, creating larger electron density above and below the molecule
26
bond length
the distance between nuclei of 2 atoms covalently bonded.
27
how does bond length affect covalent bond strength
shorter length, greater strength
28
electronegativity
the ability of an atom to attract a bonding pair of electrons.
29
how does electronegativity change over the periodic table?
decreases down groups | increases across periods
30
What are polar molecule electron densities like?
they form asymmetrical electron densities, resulting in one atom having a slight positive charge and the other a slight negative charge (delta minus and delta plus)
31
what are polar bonds represented by?
arrows from one atom to the other
32
2 bonding ideals
100% ionic | 100% covalent
33
polar bond
type of covalent bond between 2 atoms where bonding electrons are unequally distributed, meaning slight differing charges of each atom.
34
Discrete (simple molecules)
Electrically neutral group of 2 or more atoms held together by chemical bonds.
35
Dative covalent bond
Bond formed when an empty orbital of one atomoverlaps with an orbital containing a lone pair of electrons.
36
Dative bond representation
Arrow
37
Formation of Al2Cl6
Lone pair on chlorine atom overlaps with empty orbital on an aluminium atom. Only occurs in gaseous state.
38
VSEPR
Valence shell electron pair repulsion theory
39
electron pair repulsion theory basis
The shape of an electron/ion is caused by repulsion between electron pairs (lone and bond) that surround the central atom, arranging themselves to experience minimal repulsion.
40
order of repulsion strength
lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion
41
steps to work out shapes of molecules
1.) find no. of areas of electron density (bond pairs/lone pairs) via dot and cross. Then work out shape
42
2 bond pairs and no lone pair shape | angle?
linear | 180 degrees
43
3 bond pairs and no lone pairs shape and angle
trigonal planar | 120 degrees
44
4 bond pairs and no lone pairs shape | angles?
tetrahedral | 109.5 degrees
45
5 bond pairs and no lone pairs shape | angles?
trigonal bipyramidal | 90 degrees and 120 degrees
46
6 bond pairs and no lone pairs shape | angles?
hexagonal | 90 degrees and 180 degrees
47
3 bond pairs and one lone pair shape? angle? why?
trigonal pyramidal 107 degrees lone pair-lone pair repulsion is greater than bond pair-lone pair repulsion so the angles is slightly less than 109.5 degrees.
48
2 bond pairs and 2 lone pairs shape? angles?
v-shaped 104.5 degrees lone pair-lone pair repulsion > lone pair-bond pair repulsion so bond angles is further depressed.
49
dipole
a separation of charge caused by the drift of bonded electrons towards the more electronegative element.
50
why does electronegativity decrease down a group?
as atomic radii increases, distance between electrons and nuclei increases, reducing force of attraction between atoms. also more shielding
51
molecule with intermediate bonding character (both covalent and ionic)
calcium carbonate | ionic between calcium and carbonate, covalent within carbonate
52
rough size of an atom
10 to the -10 m
53
isoelectronic
having the same electronic configuration
54
size of negative ions in relation to atoms that they form from what about positive?
negative ions are larger | positive ions are smaller
55
what is ionic bonding complicated by?
ionic packing | covalent character
56
dative bonding
covalent bond in which both electrons come from the same atom
57
discrete (simple) molecule
electrically neutral group of 2 or more atoms held together by weak chemical bonds
58
DIPOLE
when 2 charges of equal magnitude but opposite signs are separated by a small distance
59
factors affecting electronegativity
nuclear charge atomic radius shielding
60
charge effects on electronegativity
higher charge, greater electronegativity
61
atomic radius effect on electronegativity
smaller atom, closer the pair of bonding atoms are to the nucleus, meaning stronger force of attraction
62
shielding effect on electronegativity
more inner shells, more shielding and so less force of attraction to attract 2 bonding pairs
63
electronegativity down a group
decreases bigger radius more shielding less force of attraction
64
electronegativity across a period
``` increases smaller radius higher charge less shielding more force of attraction ```
65
diatomic molecule polarities
non-polar | as 2 atoms in each molecule are same and so have same eneg w symmetrical densities
66
symbol representing dipole
arrow w cross through
67
linear molecules (symmetrical) polarity
non-polar as dipoles cancel each other out
68
trigonal planar molecules polarity
non-polar as is symmetrical so dipoles cancel each other out
69
tetrahedral polarity
if symmetrical, non-polar | if unsymmetrical, polar as dipoles reinforce one another
70
trigonal pyramidal polarity
dipoles reinforce one another so is polar
71
v-shaped molecule polarity
dipoles reinforce each other so is polar
72
3 types of intermolecular forces
London/dispersion forces permanent dipole interactions hydrogen bonds
73
formation of London/dispersion forces
created by instantaneous dipoles occurring due to electron density fluctuations, inducing other dipoles between other molecules, aligning so that they interact favourably with one another
74
why do London forces continue to attract each other despite movement
as fluctuations in electron density are much quicker than the kinetic energy, so continue to attract each other regardless of energy
75
factors affecting London forces
attractive force increases w increasing number of electrons in the molecule (fluctuation in density faster) more points of contact between molecules, greater overall London force
76
permanent dipole interactions
molecules possessing permanent dipoles interact and align, causing forces of attraction between molecules
77
why is London force more significant than permanent dipole interactions
random movement of molecules mean they have less alignments and so less interactions sometimes forces repel so contribute little overall
78
where do london forces exist
between all types of molecules, regardless of polarity
79
where do hydrogen bonds exist
between all atoms containing an OH group | atom bonded must be more eneg than hydrogen
80
example of hydrogen bonding
betweenoxygen atom of one molecule and hydrogen of another
81
h bonding
h bonded to an electronegative atom
82
reasons for increasing bp with increasing molecular mass
more electrons and so more instantaneous and permanent dipoles form carbon chain length increase results in more points of contact between adjacent molecules
83
branched chain effects
lower bp | due to less points of contact so decrease in IMF
84
anomalous properties of water
relatively high bp for little electrons | density of ice is less than water
85
2 conditions needing to be met for a substance to dissolve
solute particles must be surrounded by solvent particles | forces of attraction between solvent and solute particles must be strong enough to overcome solute and solute forces
86
van der waal forces
sum of al the intermolecular interactions between the molecules, including London and permanent dipole-dipole forces.
87
what happens to the ionic radius of a set of isoelectronic ions as the atomic number increases
decreases as the electrons are attracted to the nucleus more strongly.
88
giant ionic lattice
repeating units in a regular structure, caused by ions attracted in all directions to oppositely charged ions.

EDEXCEL Chemistry a level (12 decks)

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